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#1076 2021-07-10 00:20:42

ganesh
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Re: Miscellany

1054) Covalent bond

Covalent bond, in chemistry, the interatomic linkage that results from the sharing of an electron pair between two atoms. The binding arises from the electrostatic attraction of their nuclei for the same electrons. A covalent bond forms when the bonded atoms have a lower total energy than that of widely separated atoms.

A brief treatment of covalent bonds follows.

Molecules that have covalent linkages include the inorganic substances hydrogen, nitrogen, chlorine, water, and ammonia (H2, N2, Cl2, H2O, NH3) together with all organic compounds. In structural representations of molecules, covalent bonds are indicated by solid lines connecting pairs of atoms.

A single line indicates a bond between two atoms (i.e., involving one electron pair), double lines (=) indicate a double bond between two atoms (i.e., involving two electron pairs), and triple lines (≡) represent a triple bond, as found, for example, in carbon monoxide (C≡O). Single bonds consist of one sigma (σ) bond, double bonds have one σ and one pi (π) bond, and triple bonds have one σ and two π bonds.

The idea that two electrons can be shared between two atoms and serve as the link between them was first introduced in 1916 by the American chemist G.N. Lewis, who described the formation of such bonds as resulting from the tendencies of certain atoms to combine with one another in order for both to have the electronic structure of a corresponding noble-gas atom.

Covalent bonds are directional, meaning that atoms so bonded prefer specific orientations relative to one another; this in turn gives molecules definite shapes, as in the angular (bent) structure of the H2O molecule. Covalent bonds between identical atoms (as in H2) are nonpolar—i.e., electrically uniform—while those between unlike atoms are polar—i.e., one atom is slightly negatively charged and the other is slightly positively charged. This partial ionic character of covalent bonds increases with the difference in the electronegativities of the two atoms.

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It is no good to try to stop knowledge from going forward. Ignorance is never better than knowledge - Enrico Fermi. 

Nothing is better than reading and gaining more and more knowledge - Stephen William Hawking.

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#1077 2021-07-11 00:31:19

ganesh
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Registered: 2005-06-28
Posts: 33,810

Re: Miscellany

1055) Ionic bond

Ionic bond, also called electrovalent bond, type of linkage formed from the electrostatic attraction between oppositely charged ions in a chemical compound. Such a bond forms when the valence (outermost) electrons of one atom are transferred permanently to another atom. The atom that loses the electrons becomes a positively charged ion (cation), while the one that gains them becomes a negatively charged ion (anion). A brief treatment of ionic bonds follows.

Ionic bonding results in compounds known as ionic, or electrovalent, compounds, which are best exemplified by the compounds formed between nonmetals and the alkali and alkaline-earth metals. In ionic crystalline solids of this kind, the electrostatic forces of attraction between opposite charges and repulsion between similar charges orient the ions in such a manner that every positive ion becomes surrounded by negative ions and vice versa. In short, the ions are so arranged that the positive and negative charges alternate and balance one another, the overall charge of the entire substance being zero. The magnitude of the electrostatic forces in ionic crystals is considerable. Accordingly, these substances tend to be hard and nonvolatile.

An ionic bond is actually the extreme case of a polar covalent bond, the latter resulting from unequal sharing of electrons rather than complete electron transfer. Ionic bonds typically form when the difference in the electronegativities of the two atoms is great, while covalent bonds form when the electronegativities are similar.

Ionic-bond-example.jpg


It is no good to try to stop knowledge from going forward. Ignorance is never better than knowledge - Enrico Fermi. 

Nothing is better than reading and gaining more and more knowledge - Stephen William Hawking.

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#1078 2021-07-12 00:20:05

ganesh
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Re: Miscellany

1056) Hard water and Soft water

Hard water, water that contains salts of calcium and magnesium principally as bicarbonates, chlorides, and sulfates. Ferrous iron may also be present; oxidized to the ferric form, it appears as a reddish brown stain on washed fabrics and enameled surfaces. Water hardness that is caused by calcium bicarbonate is known as temporary, because boiling converts the bicarbonate to the insoluble carbonate; hardness from the other salts is called permanent. Calcium and magnesium ions in hard water react with the higher fatty acids of soap to form an insoluble gelatinous curd, thereby causing a waste of the soap. This objectionable reaction does not take place with modern detergents.

In boilers, the calcium and magnesium in hard waters form a hard, adherent scale on the plates. As a result of the poor heat conductivity of the scale, fuel consumption is increased, and the boiler deteriorates rapidly through the external overheating of the plates. Sodium carbonate, if present, hydrolyzes to produce free alkali that causes caustic embrittlement and failure of the boiler plates. Water is softened on a small scale by the addition of ammonia, borax, or trisodium phosphate, together with sodium carbonate (washing soda). The latter precipitates the calcium as carbonate and the magnesium as hydroxide. Water is softened on a large scale by the addition of just enough lime to precipitate the calcium as carbonate and the magnesium as hydroxide, whereupon sodium carbonate is added to remove the remaining calcium salts. In areas where the water is hard, home water softeners are used, making use of the properties of natural or artificial zeolite minerals. 

Soft water, water that is free from dissolved salts of such metals as calcium, iron, or magnesium, which form insoluble deposits such as appear as scale in boilers or soap curds in bathtubs and laundry equipment.

Hard water... is water that contains an appreciable quantity of dissolved minerals (like calcium and magnesium).

Soft water... is treated water in which the only ion is sodium.

Hard vs Soft

As rainwater falls, it is naturally soft. However, as water makes its way through the ground and into our waterways, it picks up minerals like chalk, lime and mostly calcium and magnesium and becomes hard water. Since hard water contains essential minerals, it is sometimes the preferred drinking water. Not only because of the health benefits, but also the flavor. On the other hand, soft water tastes salty and is sometimes not suitable for drinking. So why, then, do we soften our water?

The major difference between hard and soft water can best be seen while doing daily housework. Hard water is to blame for dingy looking clothes, dishes with spots and residue, and bathtubs with lots of film and soap scum. Soap is less effective due to its reaction to the magnesium and calcium that lather is not as rich and bubbly. Even hair washed in hard water may feel sticky and look dull. Already relating ourselves with these inconveniences? But we are not finished, hard water can take a toll on household appliances as well as use up more energy. It would also cause serious damage by scale accumulating inside water pipes, water heater, etc. These scale buildup is also known as mineral deposits or scale deposits. When water evaporates by increasing in temperature, the minerals that are suspend in precipitate causing solidified scale.

On the other hand, house workers will love using soft water, as tasks can actually be performed more efficiently with it. Soap will lather better and items will be left cleaner. Glasses will sparkle and hair will look healthy. The shower curtain will be scum-free. Clothes and skin are left softer. In addition to time, soft water can also save money, as less soap and detergents will be used. Since appliances have to work less hard, soft water can also prolong the life of washing machines, dishwaters and water heaters. Energy bills are noticeably lower when in households with water softeners. In a time when energy costs rise higher and higher, this is something for you to consider. Not to mention the benefit of reduced scale buildup.

Soft water is not, however, suggested for those with heart or circulatory problems, or others who may be on a low sodium diet. In the softening process, as minerals are removed, sodium content increases. Research shows that cardiovascular disease has the lowest risk in areas where water has the most mineral content.

There are ways to combat the sodium in soft water, which will allow households to enjoy better tasting water, as well as have the best available water for cleaning needs. They are reverse osmosis, distillation and deionization.

What type is your water? The Water Quality Association of the United States defines hard water as having dissolved mineral hardness of 1 GPG (grain per gallon) or more. Here is a helpful table to show the hardness of water:

Soft Water- less than 1 gpg
Slightly Hard- 1-3.5 gpg
Moderately Hard- 3.5-7 gpg
Very Hard- 7-10 gpg
Extremely Hard- over 10 gpg.

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It is no good to try to stop knowledge from going forward. Ignorance is never better than knowledge - Enrico Fermi. 

Nothing is better than reading and gaining more and more knowledge - Stephen William Hawking.

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#1079 2021-07-13 00:08:15

ganesh
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Re: Miscellany

1057) Fission

What is Fission?

When an atom splits into two parts, either through natural decay or when instigated within a lab, it releases energy. This process is known as fission. It has great potential as a source of power, but is also has a number of safety, environmental, and political concerns attached to it that can hinder its use.

Fission definition

An atom contains protons and neutrons in its central nucleus. In fission, the nucleus splits, either through radioactive decay or because it has been bombarded by other subatomic particles known as neutrinos. The resulting pieces have less combined mass than the original nucleus, with the missing mass converted into nuclear energy.

Controlled fission occurs when a very light neutrino bombards the nucleus of an atom, breaking it into two smaller, similarly-sized nuclei. The destruction releases a significant amount of energy — as much as 200 times that of the neutron that started the procedure — as well as releasing at least two more neutrinos.

Controlled reactions of this sort are used to release energy within nuclear power plants. Uncontrolled reactions can fuel nuclear weapons.

Radioactive fission, where the center of a heavy element spontaneously emits a charged particle as it breaks down into a smaller nucleus, does not occur often, and happens only with the heavier elements.

Fission is different from the process of fusion, when two nuclei join together rather than split apart.

Discovering atomic energy

In 1938, German physicists Otto Hahn and Fritz Strassman bombarded a uranium atom with neutrons in an attempt to make heavy elements. In a surprising twist, they wound up splitting the atom into the elements of barium and krypton, both significantly smaller than the uranium that the pair started out with. Previous efforts by physicists had resulted in only very small slivers being cut off of an atom, so the pair was puzzled by the unexpected results.

Austrian-born physicist Lise Meitner, who had fled to Sweden following Hitler's invasion of her country, realized that the split had also released energy. Working on the problem, she established that fission yielded a minimum of two neutrons for each neutron that sparked a collision. Ultimately, other physicists realized that each newly freed neutron could go on to cause two separate reactions, each of which could cause at least two more. A single impact could jumpstart a chain reaction, driving the release of still more energy.

Energy and destruction

In an intellectual chain reaction, scientists began to realize the possibilities incumbent in the new discovery. A letter to U.S. President Franklin Roosevelt at the start of World War II, drafted by Hungarian physicist Leo Szilard and signed by Albert Einstein, noted that such research could be used to create a bomb of epic proportions, and addressed the idea that the Germans could feasibly deliver such a weapon to the American doorstep. Roosevelt allocated money toward American research, and in 1941, the Office of Scientific Research and Development was formed with the aim of applying the research toward national defense.

In 1943, the Army Corp of Engineers took over the research for making a nuclear weapon. Known as the "Manhattan Project," the top-secret endeavor resulted in the formation of the first atomic bomb in July 1945. Two subsequent atomic weapons were used as part of a military strike on the cities of Hiroshima and Nagasaki in Japan.

Since then, nuclear research has been considered extremely sensitive. The knowledge itself is not overly complex, but the materials that fund the process are significantly more difficult to obtain.

More commonly, fission is used to generate energy within a nuclear power plant. However, the process creates a significant amount of nuclear waste that can be hazardous to both people and the environment. At the same time, people often fear the dangers that could come with nuclear plants and do not want them in their area. Such issues mean that nuclear energy is not as popular as more conventional methods of obtaining energy, such as the use of fossil fuels.

In the 1960s, the U.S. government explored the possibility of using fission as a method of rocket propulsion. However, the signing of the Limited (Nuclear) Test Ban Treaty in 1963 put an end to the aboveground explosion of all nuclear weapons, closing the door at least temporarily on the testing of fission-powered rockets.

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It is no good to try to stop knowledge from going forward. Ignorance is never better than knowledge - Enrico Fermi. 

Nothing is better than reading and gaining more and more knowledge - Stephen William Hawking.

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#1080 2021-07-14 01:00:20

ganesh
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Re: Miscellany

1058) Radioactivity

Radioactivity is the physical phenomenon of certain elements - such as uranium - of emitting energy in the form of radiation. This energy comes from the decay of an unstable nucleus. Any nuclear species (particular configuration of protons, neutrons and energy) that exhibit radioactivity are known as radioactive nuclei. Additionally, radioactivity or simply activity can be used as a measurement to describe how many decays a radioactive atom goes through in a period of time. These decays result in an ejection of energy and particles from the nucleus. Radioactivity can also be referred to as radioactive decay or nuclear decay.

The most common forms of radiation include alpha, beta, and gamma radiation, but other types of radioactive decay exist such as proton emission or neutron emission, or spontaneous fission of large nuclei.

Radioactivity has a number of different applications in medicine and industry. Radioactivity is even used in smoke alarms.  Additionally, these radioactive elements act as the fuel in nuclear power plants to generate electricity. As well, the radiation from these elements can be used to irradiate foods and keep them from spoiling. To learn more about uses for radioactivity of elements, see isotopes for society.

What Causes Radioactivity?

The stability of a nuclear species (also called a nuclide) is determined by forces within the nucleus. These forces determine its nuclear stability. Unstable nuclear isotopes emit radiation as a result of the conflict in the strength of the repulsive Coulomb force between protons in the nucleus and the attractive strong nuclear force between nucleons. If the Coulomb force and strong nuclear force do not balance, the nuclide in question lays outside the belt of stability and is radioactive. A number known as the neutron-to-proton ratio or N/Z ratio can be used to quickly see if the Coulomb force and strong nuclear force remain fairly balanced or out of balance. For smaller elements near the top of the periodic table, the ratio for stability is nearly 1:1. As the nuclei get larger, the N/Z ratio increases slightly for stability. If a nucleus has too many protons or neutrons, it will likely undergo some sort of transmutation to reach a more stable state (where it changes to some new nuclide with a "better" N/Z ratio).

There are many complex factors that determine whether or not a nuclide be radioactive. For example, if a nuclide has either an odd number of protons or an odd number of neutrons, it is more likely to be radioactive. If both are odd, the nuclide is almost certainly radioactive! This greater instability comes from a desire for protons and neutrons like to "pair up" with particles of the same type, boosting stability. (Of the thousands of nuclides that have been investigated only 4 stable odd-odd nuclei have been found.) Some numbers of protons or neutrons in a nucleus that promote stability. These numbers are known as magic numbers.

Most large nuclides tend to be radioactive, and the last completely stable nuclide is bismuth (which has 83 protons). These large radioactive elements often undergo alpha decay as it quickly lowers the number of protons and neutrons in the nucleus.[5] Most nuclides found in nature are not radioactive, because all of the short-lived radioactive nuclei have already decayed, leaving a vast majority of stable nuclei. There are only 50 naturally occurring nuclides that exhibit radioactivity while there are around 270 stable nuclides. Thousands of short-lived nuclides have been created in laboratories and particle accelerators.

Measuring Radioactivity

Radioactivity can also be used to describe how much ionizing radiation is released by a radioactive material. The SI unit of radioactivity is the becquerel (Bq), equal to one decay per second. The curie (Ci) was the original unit for radioactivity and is equal to 3.7×1010 Bq. Geiger counters can be used to measure the radioactivity of a substance and these devices are widely known for the "clicking" noise they make when they detect a decay producing ionizing radiation.

Another way to measure how radioactive something is is to investigate its half life, since the half life of a nuclide is related to its radiation risk.

Safety

One common misconception about radioactivity is that any radioactive object is harmful to human health. This is not the case, however, as small doses of radiation have not been proven to be harmful to humans. In fact, there are many radioactive products that can be purchased and pose no health threats to humans. Bananas, smoke detectors, some ceramic dishware, cat litter, beer, and brazil nuts are all radioactive.

However, in larger doses radiation does have negative effects on health. When radioactive materials decay, they produce ionizing radiation. Simply put, this type of radiation can strip electrons away from atoms or break chemical bonds (to make ions). This causes damage to living tissues that cannot always be repaired. Chronic exposure to radiation can lead to cancer (as a result of damage at the cellular or molecular level) or other mutations that can be harmful to fetuses. Effects from acute exposure to radiation appear quickly, and include burns and radiation poisoning. The symptoms of radiation poisoning include nausea, weakness, hair loss, and diminished organ function and this radiation sickness can result in death if the dose is high enough.

As well, some radiation is only harmful in certain circumstances. For example, smoke detectors generally contain a source of alpha particles known as americium (which is radioactive). The americium is used to detect the smoke. In the smoke detector itself, this source is radioactive but not harmful. However, because of the nature of alpha particles, the Americium is very dangerous if ingested. Smoke detectors save many lives every year, and changing the batteries once a year is a good idea. Dissecting smoke detectors however can be dangerous.
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Radioactivity refers to the particles which are emitted from nuclei as a result of nuclear instability. Because the nucleus experiences the intense conflict between the two strongest forces in nature, it should not be surprising that there are many nuclear isotopes which are unstable and emit some kind of radiation. The most common types of radiation are called alpha, beta, and gamma radiation, but there are several other varieties of radioactive decay.

Radioactive decay rates are normally stated in terms of their half-lives, and the half-life of a given nuclear species is related to its radiation risk. The different types of radioactivity lead to different decay paths which transmute the nuclei into other chemical elements. Examining the amounts of the decay products makes possible radioactive dating.

Radiation from nuclear sources is distributed equally in all directions, obeying the inverse square law.

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It is no good to try to stop knowledge from going forward. Ignorance is never better than knowledge - Enrico Fermi. 

Nothing is better than reading and gaining more and more knowledge - Stephen William Hawking.

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#1081 2021-07-15 00:30:13

ganesh
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Re: Miscellany

1059) Radium

Radium (Ra), radioactive chemical element, the heaviest of the alkaline-earth metals of Group 2 (IIa) of the periodic table. Radium is a silvery white metal that does not occur free in nature.

Element Properties

atomic number  :  88
stablest isotope  :  226
melting point  :  about 700 °C (1,300 °F)
boiling point  :  not well established (about 1,100–1,700 °C [2,000–3,100 °F])
specific gravity  :  about 5
oxidation state  :  +2

Occurrence, Properties, And Uses

Radium was discovered (1898) by Pierre Curie, Marie Curie, and an assistant, G. Bémont, after Marie Curie observed that the radioactivity of pitchblende was four or five times greater than that of the uranium it contained and was not fully explained on the basis of radioactive polonium, which she had just discovered in pitchblende residues. The new, powerfully radioactive substance could be concentrated with barium, but, because its chloride was slightly more insoluble, it could be precipitated by fractional crystallization. The separation was followed by the increase in intensity of new lines in the ultraviolet spectrum and by a steady increase in the apparent atomic weight of the material until a value of 225.2 was obtained, remarkably close to the currently accepted value of 226.03. By 1902, 0.1 gram of pure radium chloride was prepared by refining several tons of pitchblende residues, and by 1910 Marie Curie and André-Louis Debierne had isolated the metal itself.

Thirty-four isotopes of radium, all radioactive, are known; their half-lives, except for radium-226 (1,600 years) and radium-228 (5.75 years), are less than a few weeks. The long-lived radium-226 is found in nature as a result of its continuous formation from uranium-238 decay. Radium thus occurs in all uranium ores, but it is more widely distributed because it forms water-soluble compounds; Earth’s surface contains an estimated 1.8 × 10^{13} grams (2 × 10^7 tons) of radium.

Since all the isotopes of radium are radioactive and short-lived on the geological time scale, any primeval radium would have disappeared long ago. Therefore, radium occurs naturally only as a disintegration product in the three natural radioactive decay series (thorium, uranium, and actinium series). Radium-226 is a member of the uranium-decay series. Its parent is thorium-230 and its daughter radon-222. The further decay products, formerly called radium A, B, C, C′, C″, D, and so on, are isotopes of polonium, lead, bismuth, and thallium.

Compounds

The chemistry of radium is what would be expected of the heaviest of the alkaline earths, but the intense radioactivity is its most characteristic property. Its compounds display a faint bluish glow in the dark, a result of their radioactivity in which emitted alpha particles excite electrons in the other elements in the compound and the electrons release their energy as light when they are de-excited. One gram of radium-226 undergoes 3.7 × 10^{10} disintegrations per second, a level of activity that defined the curie (Ci), an early unit of radioactivity. This is an energy release equivalent to about 6.8 × 10^{−3} calorie per second, sufficient to raise the temperature of a well-insulated 25-gram sample of water at the rate of 1 °C every hour. The practical energy release is even greater than this (by four to five times), because of the production of a large number of short-lived radioactive decay products. The alpha particles emitted by radium may be used to initiate nuclear reactions.

Radium’s uses all stem from its radioactivity. The most important use of radium was formerly in medicine, principally for the treatment of cancer by subjecting tumours to the gamma radiation of its daughter isotopes. Radium-223, an alpha emitter with a half-life of 11.43 days, has been studied for use in cell-directed cancer therapy, in which a monoclonal antibody or related targeting protein with high specificity is attached to the radium. In most therapeutic applications, however, radium has been superseded by the less costly and more powerful artificial radioisotopes cobalt-60 and cesium-137. An intimate mixture of radium and beryllium is a moderately intense source of neutrons and has been used for scientific research and for well logging in geophysical prospecting for petroleum. For these uses, however, substitutes have become available. One of the products of radium decay is radon, the heaviest noble gas; this decay process is the chief source of that element. A gram of radium-226 will emit 1 × 10^{−4} millilitre of radon per day.

When a radium salt is mixed with a paste of zinc sulfide, the alpha radiation causes the zinc sulfide to glow, yielding a self-luminescent paint for watch, clock, and instrument dials. From about 1913 up until the 1970s, several million radium dials, coated with a mixture of radium-226 and zinc sulfide, were manufactured. By the early 1930s it was found, however, that exposure to radium posed a serious hazard to health: a number of women who had worked with the radium-containing luminescent paint during the 1910s and ’20s subsequently died. They had ingested considerable amounts of radium through the technique called “lip-pointing,” which meant using their lips and tongues to shape their paintbrushes to a fine tip. Like calcium and strontium, radium tends to concentrate in bones, where its alpha radiation interferes with red corpuscle production, and some of those women developed anemia and bone cancer. The practice of employing radium in luminescent coatings was curtailed in the early 1960s after the high toxicity of the material was recognized. Phosphorescent paints that absorb light and later release it have replaced radium. (The detection of exhaled radon provides a very sensitive test for radium absorption.)

Radium metal may be prepared by electrolytic reduction of its salts, and it displays high chemical reactivity. It is attacked by water with vigorous evolution of hydrogen and by air with the formation of the nitride. It occurs exclusively as the Ra2+ ion in all its compounds. The sulfate, RaSO4, is the most insoluble sulfate known, and the hydroxide, Ra(OH)2, is the most soluble of the alkaline-earth hydroxides. The gradual buildup of helium within crystals of radium bromide, RaBr2, weakens them, and they occasionally explode. In general, the compounds of radium are very similar to their barium counterparts, making separation of the two elements difficult.

In modern technology, radium is separated from barium by fractional crystallization of the bromides, followed by purification through ion-exchange techniques for removal of the last 10 percent of the barium.

radium.jpg


It is no good to try to stop knowledge from going forward. Ignorance is never better than knowledge - Enrico Fermi. 

Nothing is better than reading and gaining more and more knowledge - Stephen William Hawking.

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#1082 2021-07-16 00:44:51

ganesh
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Re: Miscellany

1060) Actinium

Actinium (Ac), radioactive chemical element, in Group 3 (IIIb) of the periodic table, atomic number 89. Actinium was discovered (1899) by French chemist André-Louis Debierne in pitchblende residues left after French physicists Pierre and Marie Curie had extracted radium from them, and it was also discovered (1902) independently by German chemist Friedrich Oskar Giesel. Debierne named the element after the Greek word aktinos (“ray”). A ton of pitchblende ore contains about 0.15 mg of actinium. The rare silvery-white metal is highly radioactive, glowing blue in the dark.

The most common isotope of actinium is actinium-227; the others, natural and artificial, are too short-lived to accumulate in macroscopic quantity. Actinium-227, which is one of the decay products of uranium-235, has a 21.8-year half-life and in turn decays almost entirely to thorium-227, but about 1 percent decays to francium-223. This whole disintegration chain with its branches is called the actinium series.

Actinium-225 has a 10-day half-life, decaying by the emission of alpha particles. Its short-lived daughter isotopes emit only alpha and beta particles with no high-energy gamma rays. This isotope can thus deliver high-energy radiation to a tumour without greatly affecting the surrounding tissue. Complexes of actinium-225 have been studied for their use in nuclear medicine.

Actinium, the ions of which in solution are colourless, exhibits an oxidation state of +3, closely resembling the rare-earth lanthanoid elements in its chemical properties. Actinium is the prototype of a second rare-earth-like series, the actinoid elements.

Element Properties

atomic number  :  89
stablest isotope  :  227
oxidation state  :  +3
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actiniumcolor_p.png


It is no good to try to stop knowledge from going forward. Ignorance is never better than knowledge - Enrico Fermi. 

Nothing is better than reading and gaining more and more knowledge - Stephen William Hawking.

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#1083 2021-07-17 00:31:22

ganesh
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Re: Miscellany

1061) Molybdenum

Molybdenum (Mo), chemical element, silver-gray refractory metal of Group 6 (VIb) of the periodic table, used to impart superior strength to steel and other alloys at high temperature.

The Swedish chemist Carl Wilhelm Scheele had demonstrated (c. 1778) that the mineral molybdaina (now molybdenite), for a long time thought to be a lead ore or graphite, certainly contains sulfur and possibly a previously unknown metal. At Scheele’s suggestion, Peter Jacob Hjelm, another Swedish chemist, successfully isolated the metal (1782) and named it molybdenum, from the Greek molybdos, “lead.”

Molybdenum is not found free in nature. A relatively rare element, it is about as abundant as tungsten, which it resembles. For molybdenum the chief ore is molybdenite—molybdenum disulfide, MoS2—but molybdates such as lead molybdate, PbMoO4 (wulfenite), and MgMoO4 are also found. Most commercial production is from ores containing the mineral molybdenite. The concentrated mineral is usually roasted in an excess of air to yield molybdenum trioxide (MoO3), also called technical molybdic oxide, which, after purification, can be reduced with hydrogen to the metal. Subsequent treatment depends on the ultimate use of molybdenum. Molybdenum may be added to steel in the furnace in the form of either technical oxide or ferromolybdenum. Ferromolybdenum (containing at least 60 percent molybdenum) is produced by igniting a mixture of technical oxide and iron oxide. Molybdenum metal is produced in the form of a powder by hydrogen reduction of chemically pure molybdic oxide or ammonium molybdate, (NH4)2MoO4. The powder is converted to massive metal by the powder-metallurgy process or by the arc-casting process.

Molybdenum-base alloys and the metal itself have useful strength at temperatures above which most other metals and alloys are molten. The major use of molybdenum, however, is as an alloying agent in the production of ferrous and nonferrous alloys, to which it uniquely contributes hot strength and corrosion resistance, e.g., in jet engines, combustion liners, and afterburner parts. It is one of the most effective elements for increasing hardenability of iron and steel, and it also contributes to the toughness of quenched and tempered steels. The high corrosion resistance needed in the stainless steels used for processing pharmaceuticals and in the chromium steels for automotive trim is uniquely enhanced by small additions of molybdenum. Metallic molybdenum has been used for such electric and electronic parts as filament supports, anodes, and grids. Rod or wire is used for heating elements in electric furnaces operating up to 1,700 °C (3,092 °F). Coatings of molybdenum adhere firmly to steel, iron, aluminum, and other metals and show excellent resistance to wear.

Molybdenum is rather resistant to attack by acids, except for mixtures of concentrated nitric and hydrofluoric acids, and it can be attacked rapidly by alkaline oxidizing melts, such as fused mixtures of potassium nitrate and sodium hydroxide or sodium peroxide; aqueous alkalies, however, are without effect. It is inert to oxygen at normal temperature but combines with it readily at red heat, to give the trioxides, and is attacked by fluorine at room temperature, to give the hexafluorides.

Natural molybdenum is a mixture of seven stable isotopes: molybdenum-92 (15.84 percent), molybdenum-94 (9.04 percent), molybdenum-95 (15.72 percent), molybdenum-96 (16.53 percent), molybdenum-97 (9.46 percent), molybdenum-98 (23.78 percent), and molybdenum-100 (9.13 percent). Molybdenum exhibits oxidation states of +2 to +6 and is considered to display the zero oxidation state in the carbonyl Mo(CO)6. Molybdenum(+6) appears in the trioxide, the most important compound, from which most of its other compounds are prepared, and in the molybdates, used to produce pigments and dyes. Molybdenum disulfide (MoS2), which resembles graphite, is used as a solid lubricant or as an additive to greases and oils. Molybdenum forms hard, refractory, and chemically inert interstitial compounds with boron, carbon, nitrogen, and silicon upon direct reaction with those elements at high temperatures.

Molybdenum is an essential trace element in plants; in legumes as a catalyst it assists bacteria in fixing nitrogen. Molybdenum trioxide and sodium molybdate (Na2MoO4) have been used as micronutrients.

The largest producers of molybdenum are China, the United States, Chile, Peru, Mexico, and Canada.

Element Properties

atomic number  :  42
atomic weight  :  95.94
melting point  :  2,610 °C (4,730 °F)
boiling point  :  5,560 °C (10,040 °F)
specific gravity  :  10.2 at 20 °C (68 °F)
oxidation states  :  0, +2, +3, +4, +5, +6

Mo.jpg


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#1084 2021-07-18 00:24:29

ganesh
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Re: Miscellany

1062) Astatine

Astatine (At), radioactive chemical element and the heaviest member of the halogen elements, or Group 17 (VIIa) of the periodic table. Astatine, which has no stable isotopes, was first synthetically produced (1940) at the University of California by American physicists Dale R. Corson, Kenneth R. MacKenzie, and Emilio Segrè, who bombarded bismuth with accelerated alpha particles (helium nuclei) to yield astatine-211 and neutrons. Naturally occurring astatine isotopes have subsequently been found in minute amounts in the three natural radioactive decay series, in which they occur by minor branching (astatine-218 in the uranium series, astatine-216 in the thorium series, and astatine-215 and astatine-219 in the actinium series). Thirty-two isotopes are known; astatine-210, with a half-life of 8.1 hours, is the longest lived. Because astatine has no stable or long-lived isotopes, it was given its name from the Greek word astatos, meaning “unstable.”

Element Properties

atomic number  :  85
stablest isotope  :  210
oxidation states      :  −1, +1, +3(?), +5, +7(?)

Production And Use

The only practical way of obtaining astatine is by synthesizing it through nuclear reactions.

It indicates that bismuth-209 takes up one alpha particle and emits x neutrons to form an isotope of astatine, whose atomic weight depends on the number of neutrons lost. Metallic bismuth may be used as a target material. From this, astatine may readily be removed by distillation in air from a stainless-steel tube. The free element begins to distill at 271 °C (520 °F, or the melting point of bismuth), but the operation is best carried out at 800 °C (1,500 °F) with subsequent redistillation. If an aqueous solution of astatine is desired, the element may be separated by washing with an appropriate aqueous solution. Alternatively, the halogen may be removed from the target by chemical methods, such as dissolving in nitric acid, with the latter being removed by boiling.

Another procedure involves the use of a metallic thorium target, which—after bombardment—is dissolved in concentrated hydrochloric acid containing hydrogen fluoride and chlorine.

Analysis

Because of the short half-lives of astatine isotopes, only very small quantities have been available for study. With the exception of a few spectrometric and mass-spectrometric studies, most investigations of astatine chemistry have used tracer techniques, which involve using chemical reactions in a solution with similarly reacting elements as carriers. The amount of astatine is then calculated from the measured radioactivity of the reaction products. However, the rarity of astatine means that these solutions are extremely dilute, with concentrations around or below {10}^{−10} molarity (the number of moles per litre of solution). At such concentrations, the effects of impurities can be very serious, especially for a halogen such as astatine, which exists in several oxidation states and can form many organic compounds. Iodine has been used as a carrier in most experiments. Techniques applied include coprecipitation, solvent extraction, ion exchange, and other forms of chromatography (separation by adsorption differences), electrodeposition (deposition by an electric current), electromigration (movement in an electric field), and diffusion. A direct identification of some astatine compounds has been made by mass spectrometry.

Except for nuclear properties, the only physical property of astatine to be measured directly is the spectrum of atomic astatine. Other physical properties have been predicted from theory and by extrapolation from the properties of other elements.

Chemical Properties

Some of the chemical properties of the element have been established. It generally resembles iodine. Thus, like iodine, it concentrates in the thyroid gland of higher animals. A substantial portion, however, is distributed throughout the body and acts as an internal radiation source.

The astatide ion, At−, is quantitatively coprecipitated with insoluble iodides, such as silver iodide or thallium iodide. The diffusion coefficient of the iodide ion is 1.42 times that of the astatide ion, which moves more slowly toward the anode than the former under given conditions. The ion is formed by reduction of the element, using zinc or sulfur dioxide. It is oxidized to the zero valence state by the ferric ion, Fe3+, iodine (I2), and dilute nitric acid. Thus, the astatide ion is a stronger reducing agent than the iodide ion, and free iodine is a stronger oxidizing agent than astatine.

Free astatine is characterized by volatility from solution and by extractability into organic solvents. It undergoes disproportionation in alkaline media. Astatine is coprecipitated with cesium iodide and thus appears to form polyhalide anions. Astatine extracted into chloroform has been shown to coprecipitate homogeneously with iodine when a portion of the latter is crystallized. Astatine seems to be present as the iodide, which appears to be more polar (i.e., showing separation of electric charge) in character than iodine bromide. It is somewhat soluble in water and much more soluble in benzene and carbon tetrachloride.

Astatine is known to occur in positive oxidation numbers. The astatate ion, (AtO3)−, is coprecipitated with insoluble iodates, such as silver iodate (AgIO3), and is obtained by the oxidation of lower oxidation states with hypochlorite, periodate, or persulfate. So far no evidence for perastatate has been found, but this may be because the ion, (AtO6)5−, may show little tendency to coprecipitate with potassium iodate (KIO4).

Astatine in the +1 state is stabilized by complexation, and complexes formulated as dipyridine astatine perchlorate [At(py)2] [ClO4] and dipyridine astatine nitrate [At(py)2] [NO3] have been prepared. Compounds with the formulas (C6H5)AtCl2, (C6H5)2AtCl, and (C6H5)AtO2 have also been obtained. A variety of methods may be used to synthesize astatobenzene, C6H5At.
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#1085 2021-07-19 00:08:23

ganesh
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Re: Miscellany

1063) Rubidium

Rubidium (Rb), chemical element of Group 1 (Ia) in the periodic table, the alkali metal group. Rubidium is the second most reactive metal and is very soft, with a silvery-white lustre.

Rubidium was discovered (1861) spectroscopically by German scientists Robert Bunsen and Gustav Kirchhoff and named after the two prominent red lines of its spectrum. Rubidium and cesium often occur together in nature. Rubidium, however, is more widely scattered and seldom forms a natural mineral; it is found only as an impurity in other minerals, ranging in content up to 5 percent in such minerals as lepidolite, pollucite, and carnallite. Brine samples have also been analyzed that contain up to 6 parts per million of rubidium.

In the principal commercial process of rubidium production, small amounts of rubidium are obtained from the mixture of alkali metal carbonates remaining after lithium salts are extracted from lepidolite. Primarily a potassium carbonate, this by-product also contains approximately 23 percent rubidium and 3 percent cesium carbonates.

The primary difficulty associated with the production of pure rubidium is that it is always found together with cesium in nature and is also mixed with other alkali metals. Because these elements are very similar chemically, their separation presented numerous problems before the advent of ion-exchange methods and ion-specific complexing agents such as crown ethers. Once pure salts have been prepared, it is a straightforward task to convert them to the free metal. This can be done by electrolysis of the fused cyanide or by reduction with calcium or sodium followed by fractional distillation.

Rubidium is difficult to handle because it ignites spontaneously in air, and it reacts violently with water to yield a solution of rubidium hydroxide (RbOH) and hydrogen, which bursts into flames; rubidium is therefore kept in dry mineral oil or an atmosphere of hydrogen. If a metal sample has a large enough surface area, it can burn to form superoxides. Rubidium superoxide (RbO2) is a yellow powder. Rubidium peroxides (Rb2O2) can be formed by oxidation of the metal with the required amount of oxygen. Rubidium forms two other oxides (Rb2O and Rb2O3).

It is used in photoelectric cells and as a “getter” in electron tubes to scavenge the traces of sealed-in gases. Rubidium atomic clocks, or frequency standards, have been constructed, but they are not as precise as cesium atomic clocks. However, aside from these applications, rubidium metal has few commercial uses and is of very minor economic significance. High prices and an uncertain and limited supply discourage the development of commercial uses.

Natural rubidium makes up about 0.01 percent of Earth’s crust; it exists as a mixture of two isotopes: rubidium-85 (72.15 percent) and the radioactive rubidium-87 (27.85 percent), which emits beta rays with a half-life of about 6 × 1011 years. A large number of radioactive isotopes have been artificially prepared, from rubidium-79 to rubidium-95. One estimate of the age of the solar system as 4.6 billion years is based on the ratio of rubidium-87 to strontium-87 in a stony meteorite. Rubidium easily loses its single valence electron but no others, accounting for its oxidation number of +1, although several compounds that contain the anion, Rb-, have been synthesized.

Rubidium and cesium are miscible in all proportions and have complete solid solubility; a melting-point minimum of 9 °C (48 °F) is reached. Rubidium forms a number of mercury amalgams. Because of the increased specific volume of rubidium, as compared with the lighter alkali metals, there is a lesser tendency for it to form alloy systems with other metals.

Element Properties

atomic number  :  37
atomic weight  :  85.47
melting point  :  38.9 °C (102 °F)
boiling point  :  688 °C (1,270 °F)
specific gravity  :  1.53 (at 20 °C, or 68 °F)
oxidation states  :  +1, -1 (rare)
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Rubidium-Pictures.jpg


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#1086 2021-07-20 00:23:29

ganesh
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Re: Miscellany

1064) Benzoic acid

Benzoic acid, a white, crystalline organic compound belonging to the family of carboxylic acids, widely used as a food preservative and in the manufacture of various cosmetics, dyes, plastics, and insect repellents.

First described in the 16th century, benzoic acid exists in many plants; it makes up about 20 percent of gum benzoin, a vegetable resin. It was first prepared synthetically about 1860 from compounds derived from coal tar. It is commercially manufactured by the chemical reaction of toluene (a hydrocarbon obtained from petroleum) with oxygen at temperatures around 200° C (about 400° F) in the presence of cobalt and manganese salts as catalysts. Pure benzoic acid melts at 122° C (252° F) and is very slightly soluble in water.

Among the derivatives of benzoic acid are sodium benzoate, a salt used as a food preservative; benzyl benzoate, an ester used as a miticide; and benzoyl peroxide, used in bleaching flour and in initiating chemical reactions for preparing certain plastics.

b1676.png
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#1087 2021-07-21 00:28:24

ganesh
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Re: Miscellany

1065) Tartaric acid

Tartaric acid, also called dihydroxybutanedioic acid, a dicarboxylic acid, one of the most widely distributed of plant acids, with a number of food and industrial uses. Along with several of its salts, cream of tartar (potassium hydrogen tartrate) and Rochelle salt (potassium sodium tartrate), it is obtained from by-products of wine fermentation. In a partially purified form, tartar was known to the ancient Greeks and Romans; the free acid was first isolated in 1769 by Swedish chemist Carl Wilhelm Scheele. The lees, or sediments, and other waste products from fermentation are heated and neutralized with calcium hydroxide; the precipitated calcium tartrate is then treated with sulfuric acid to produce free tartaric acid. Rochelle salt is prepared from the crude crystalline potassium acid salt, called argol, by neutralization with sodium carbonate. Purified cream of tartar comes chiefly from the filtrates from production of the acid and Rochelle salt. A third salt, tartar emetic (antimony potassium tartrate), is made from the potassium acid salt and antimony oxide.

Three stereoisomeric forms of tartaric acid exist: (1) dextrorotatory tartaric acid (D-tartaric acid) found in grapes and several other fruits, (2) levorotatory tartaric acid (L-tartaric acid) obtained chiefly by resolution of racemic tartaric acid, and (3) a meso or achiral form. Racemic tartaric acid (an equal mixture of D- and L-tartaric acid) is prepared commercially by the molybdenum- or tungsten-catalyzed oxidation of maleic anhydride with hydrogen peroxide.

Study of the crystallographic, chemical, and optical properties of the tartaric acids by French chemist and microbiologist Louis Pasteur laid the basis for modern ideas of stereoisomerism.

The various tartaric acids and the common tartrate salts are all colourless, crystalline solids readily soluble in water. Tartaric acid is widely used as an acidulant in carbonated drinks, effervescent tablets, gelatin desserts, and fruit jellies. It has many industrial applications—e.g., in cleaning and polishing metals, in calico printing, in wool dyeing, and in certain photographic printing and development processes. Rochelle salt is used in silvering mirrors, in processing cheese, and in compounding mild cathartics. Cream of tartar is incorporated into baking powders, hard candies, and taffies; and it is employed in the cleaning of brass, the electrolytic tinning of iron and steel, and the coating of other metals with gold and silver. Tartar emetic is used as an insecticide and a dyeing mordant.

1-3-The-structural-formula-of-L-tartaric-acid-Ill-by-Ida-S-R-Stenhagen.png


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#1088 2021-07-22 00:45:11

ganesh
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Re: Miscellany

1066) Lactic acid

Lactic acid, also called α-hydroxypropionic acid, or 2-hydroxypropanoic acid, an organic compound belonging to the family of carboxylic acids, present in certain plant juices, in the blood and muscles of animals, and in the soil. It is the commonest acidic constituent of fermented milk products such as sour milk, cheese, and buttermilk.

First isolated in 1780 by a Swedish chemist, Carl Wilhelm Scheele, lactic acid is manufactured by the fermentation of molasses, starch, or whey in the presence of alkaline substances such as lime or calcium carbonate; it is available as aqueous solutions of various concentrations, usually 22–85 percent, and degrees of purity. Lactic acid is used in tanning leather and dyeing wool; as a flavouring agent and preservative in processed cheese, salad dressings, pickles, and carbonated beverages; and as a raw material or a catalyst in numerous chemical processes. Pure lactic acid, rarely prepared, is a colourless, crystalline substance that melts at 18° C (64° F); it rapidly absorbs moisture from the atmosphere.

Lactic acid occurs in the blood (in the form of its salts, called lactates) when glycogen is broken down in muscle and can be converted back to glycogen in the liver. Lactates are also the products of fermentation (q.v.) in certain bacteria..

Lactic_acid_molecule.png


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#1089 Yesterday 00:13:01

ganesh
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Re: Miscellany

1067) Bromine

Bromine (Br), chemical element, a deep red noxious liquid, and a member of the halogen elements, or Group 17 (Group VIIa) of the periodic table.

Element Properties

atomic number  :  35
atomic weight  :  79.909
melting point  :  −7.2 °C (19 °F)
boiling point  :  59 °C (138 °F)
specific gravity  :  3.12 at 20 °C (68 °F)
oxidation states  :  −1, +1, +3, +5, +7

History

Bromine was discovered in 1826 by the French chemist Antoine-Jérôme Balard in the residues (bitterns) from the manufacture of sea salt at Montpellier. He liberated the element by passing chlorine through an aqueous solution of the residues, which contained magnesium bromide. Distillation of the material with manganese dioxide and sulfuric acid produced red vapours, which condensed to a dark liquid. The similarity of this procedure to that for making chlorine suggested to Balard that he had obtained a new element similar to chlorine. (The German chemist Justus von Liebig appears to have obtained the element before Balard, but he wrongly considered it to be iodine chloride.) Because of the bad odour of the element, the French Academy of Sciences suggested the name bromine, from the Greek word bromos, meaning “bad smell” or “stench.”

Occurrence And Distribution

A rare element, bromine is found in nature dispersed throughout Earth’s crust only in compounds as soluble and insoluble bromides. Some enrichment occurs in ocean water (65 parts per million by weight), in the Dead Sea (approximately 5 grams per litre [0.7 ounce per gallon]), in some thermal springs, and in rare insoluble silver bromide minerals (such as bromyrite, found in Mexico and Chile). Natural salt deposits and brines are the main sources of bromine and its compounds. Jordan, Israel, China, and the United States led the world in bromine production in the early 21st century; other important bromine-producing countries during that period include Japan, Ukraine, and India.

Natural bromine is a mixture of two stable isotopes: bromine-79 (50.54 percent) and bromine-81 (49.46 percent). Of the 17 known radioactive isotopes of the element, bromine-77 has the longest half-life (57 hours).

Physical And Chemical Properties

Free bromine is a reddish brown liquid with an appreciable vapour pressure at room temperature. Bromine vapour is amber in colour. Bromine has a pungent odour and is irritating to the skin, eyes, and respiratory system. Exposure to concentrated bromine vapour, even for a short time, may be fatal. Like the other halogens, bromine exists as diatomic molecules in all aggregation states.

About 3.41 grams (0.12 ounce) of bromine dissolve in 100 millilitres (0.1 quart) of water at room temperature. The solution is known as bromine water. Like chlorine water, it is a good oxidizing agent, and it is more useful because it does not decompose so readily. It liberates free iodine from iodide-containing solutions and sulfur from hydrogen sulfide. Sulfurous acid is oxidized by bromine water to sulfuric acid. In sunlight bromine water decomposes, with release of oxygen.

From bromine water a hydrate (a clathrate) can be isolated that contains 172 water molecules and 20 cavities capable of accommodating the bromine molecules. Bromine dissolves in aqueous alkali hydroxide solutions, giving bromides, hypobromites, or bromates, depending on the temperature. Bromine is readily extracted from water by organic solvents such as carbon tetrachloride, chloroform, or carbon disulfide, in which it is very soluble. In the organic solvents it gives an orange solution.

The electron affinity of bromine is high and is similar to that of chlorine. It is, however, a less powerful oxidizing agent, chiefly because of the weaker hydration of the bromide ion as compared with the chloride ion. Similarly, a metal-bromine bond is weaker than the corresponding metal-chlorine bond, and this difference is reflected in the chemical reactivity of bromine, which lies between that of chlorine and that of iodine. An organic bromo compound resembles the corresponding chloro derivative but is usually more dense, less volatile, less combustible, and less stable.

Bromine combines violently with the alkali metals and with phosphorus, As, aluminum, and antimony but less violently with certain other metals. Bromine displaces hydrogen from saturated hydrocarbons and adds to unsaturated hydrocarbons, though not as readily as chlorine does.

The most stable oxidation state of the element is −1, in which bromine occurs naturally. But oxidation states of 0 (elemental bromine, Br2), +1 (hypobromite, BrO−), +3 (bromite, BrO−2), +5 (bromate, BrO−3), and +7 (perbromate, BrO−4) are also known. The first ionization energy of bromine is high, and compounds containing bromine in positive oxidation numbers are stabilized by appropriate ligands, mainly oxygen and fluorine. Compounds with the oxidation numbers +1, +3, +4, +5, and +7 all contain covalent bonds.

Production And Use

The chief commercial source of bromine is ocean water, from which the element is extracted by means of chemical displacement (oxidation) by chlorine in the presence of sulfuric acid.

The product of the reaction is a dilute solution of bromine, from which the element is removed by blowing air through it. The free bromine is then mixed with sulfur dioxide, and the mixed gases are passed up a tower down which water is trickling.

Resulting in a mixture of acids that is much richer in bromide ion than seawater. A second treatment with chlorine liberates bromine, which is freed from chlorine and purified by passage over moist iron filings.

Commercial bromine generally contains up to 0.3 percent chlorine. It is usually stored in glass bottles or in barrels coated with lead or Monel metal.

The industrial usage of bromine had been dominated by the compound ethylene bromide (C2H4Br2), which once was added to gasoline with tetraethyl lead to prevent deposition of lead in the engine. Since the renunciation of leaded gasoline, bromine compounds have mainly been used in flame retardants, but ethylene bromide is still an important compound because of its use to destroy nematodes and other pests in soils. Bromine is also used in the production of catalysts, such as aluminum bromide.

Bromine has other uses, as in making various dyes and the compounds tetrabromoethane (C2H2Br4) and bromoform (CHBr3), which are used as liquids in gauges because of their high specific gravity. Until the development of barbiturates in the early 20th century, bromides of potassium, sodium, calcium, strontium, lithium, and ammonium were used widely in medicine because of their sedative action. Silver bromide (AgBr), an important component of photographic film, is, like silver chloride and iodide, light sensitive. Traces of potassium bromate (KBrO3) are added to wheat flour to improve baking. Other bromine compounds of significance include hydrogen bromide (HBr), a colourless gas used as a reducing agent and a catalyst in organic reactions. A solution of the gas in water is called hydrobromic acid, a strong acid that resembles hydrochloric acid in its activity toward metals and their oxides and hydroxides.

Analysis

A sensitive test for bromine is the reaction with fluorescein to give a deep red colour caused by bromination of the organic molecule, or by its reaction with fuchsine dyes in the presence of sulfurous acid, to give a deep blue colour. A more common test involves heating the sample with dilute sulfuric acid in the presence of potassium dichromate; the bromine is then extracted with chloroform, and, upon addition of potassium iodide, the pink colour of iodine appears. The presence of bromine may also be recognized by the evolution of hydrogen bromide containing some brown bromine vapour when a solid sample is treated with concentrated sulfuric acid. Alternatively, chlorine may be added to an aqueous solution of a sample containing bromide, causing development of a brown colour (free bromine).

For the quantitative determination of bromine, the following methods are recommended:
1.    Free bromine is titrated with sodium thiosulfate in the presence of potassium iodide.
2.    Bromides may be determined either gravimetrically (by weight analysis) or by titration with silver nitrate.
3.    In the presence of chloride and iodide, the potentiometric method may be used (as with chlorine).
4.    In the absence of iodide, bromide may be oxidized to bromine, which is then determined in the distillate. Alternatively, bromide may be oxidized to bromate by hypochlorous acid. The excess of the oxidizing agent is destroyed by sodium formate, and iodine is liberated by addition of potassium iodide and acid, with the free iodine being titrated by thiosulfate.
5.    For the determination of bromine in an organic compound, the latter is oxidized by nitric acid, and the bromine is determined as silver bromide.

Bromine.jpg


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#1090 Today 00:14:13

ganesh
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Re: Miscellany

1068) Virology

Virology, branch of microbiology that deals with the study of viruses.

Although diseases caused by viruses have been known since the 1700s and cures for many were (somewhat later) effected, the causative agent was not closely examined until 1892, when a Russian bacteriologist, D. Ivanovski, observed that the causative agent (later proved to be a virus) of tobacco mosaic disease could pass through a porcelain filter impermeable to bacteria. Modern virology began when two bacteriologists, Frederick William Twort in 1915 and Félix d’Hérelle in 1917, independently discovered the existence of bacteriophages (viruses that infect bacteria).

Direct visualization of viruses became possible after the electron microscope was introduced about 1940. In 1935 tobacco mosaic virus became the first virus to be crystallized; in 1955 the poliomyelitis virus was crystallized. (A virus “crystal” consists of several thousand viruses and, because of its purity, is well suited for chemical studies.) Virology is a discipline of immediate interest because many human diseases, including smallpox, influenza, the common cold, and AIDS, as well as a host of economically important plant and animal diseases, are caused by viruses.

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